Transition metals: Some chemistry of iron (2)

Some chemistry of Iron (2):

Edexcel

24. be able to record observations and write suitable equations for the reactions of Cr³⁺(aq), Fe²⁺(aq), Fe³⁺(aq), Co²⁺(aq) and Cu²⁺(aq) with aqueous sodium hydroxide and aqueous ammonia, including in excess.

25. be able to write ionic equations to show the difference between ligand exchange and amphoteric behaviour for the reactions in (24) above.

34. understand the role of Fe²⁺ ions in catalysing the reaction between I− and S₂O₈^2— ions.

Manganate(VII) with iron (II) titration self-indicating

catalyst in Haber process

AQA

Exchange of the ligand H₂O by Cl– can involve a change of co-ordination number (e.g. Fe³⁺(aq), Co²⁺(aq) and Cu²⁺(aq).

Haem is an iron(II) complex with a multidentate ligand.

Oxygen forms a co-ordinate bond to Fe(II) in haemoglobin, enabling oxygen to be transported in the blood.

Carbon monoxide is toxic because it replaces oxygen co-ordinately bonded to Fe(II) in haemoglobin.

The redox titrations of Fe²⁺(aq) and C₂O₄^2– with MnO₄^—

Students should be able to perform calculations for these titrations and similar redox reactions.

Examples include, finding:

•   the mass of iron in an iron tablet 


•   the percentage of iron in steel 


•   the Mr of hydrated ammonium iron(II) sulfate 


•       Fe is used as a heterogeneous catalyst in the Haber process.

In aqueous solution, the following metal-aqua ions are formed:

[M(H₂O)₆]²⁺, limited to M = Fe and Cu

[M(H₂O)₆]³⁺, limited to M = Al and Fe

The acidity of [M(H₂O)₆]³⁺ is greater than that of [M(H₂O)₆]²⁺

Students should be able to:

•   explain, in terms of the charge/size ratio of the metal ion, why the acidity of [M(H₂O)₆]³⁺ is greater than that of [M(H₂O)₆]²⁺ 


•   describe and explain the simple test-tube reactions of M²⁺(aq) ions, limited to M = Fe and Cu, and of M³⁺(aq) ions, limited to M = Al and Fe, with the bases OH–, NH3 and CO₃^2— 


•       Students could carry out test-tube reactions of metal-aqua ions with NaOH, NH₃ and Na₂CO₃

Halogen carrier in arene organic chemistry

OCR

Redox reactions


(k) redox reactions and accompanying colour changes for:

(i) interconversions between Fe2+ and Fe3+

Fe2+ can be oxidised with H+/MnO4– and Fe3+ reduced with I–

7. Iron as a catalyst

Haber Process

a) The Haber process is used to form ammonia from atmospheric nitrogen and hydrogen extracted from methane (natural gas). 

The industrial process uses iron as a catalyst in a process in which the pressure is around 200atm and the working temperature around 450^oC. 

N₂(g)    +     3H₂(g)     ⇌     2NH₃(g)

You can find out more about the Haber Process if you follow this link:

https://derekcarrsavvy-chemist.blogspot.co.uk/2016/10/equilibrium-i-applying-le-chateliers.html

Iron here acts as a heterogenous catalyst because it is in a different state to the gaseous reactants.

b) Redox catalysis of the peroxodisulphate/iodide reaction

The reaction between peroxodisulphate ions and iodide ions is typically used to illustrate the action of iron ions and their catalytic effect in this reaction. 

This is the equation for the redox reaction between peroxodisulphate ions and iodide ions. 

2I^—      +      S₂O₈^2—          ⟶        I₂      +      2SO₄^2—

As you can see both ions are anions negatively charged and so will repel each other.

The result is that this reaction has a very high activation energy. 

Iron ions have a catalytic effect because they lower the activation energy of the reaction. 

When a small amount of iron(II) ions (Fe²⁺) are added to the iodide -peroxodisulphate mixture, the peroxodisulphate ions oxidise them to iron(III) ions. 

2Fe²⁺        +      S₂O₈^2—          ⟶        2Fe³⁺      +      2SO₄^2—

Then the iron(III) ions formed oxidise the iodide ions to iodine and they are reduced back to iron(II)

2I^—      +      2Fe³⁺           ⟶        I₂      +      2Fe²⁺       

These iron(II) ions can now start the redox process all over again.

The reaction is much faster because the two reactions catalysed with iron ions have much lower activation energies. 

In this example the iron(II) ions act as a homogenous catalyst since they are in the same state as the iodide and peroxodisulphate ions. 

8. Titration of iron(II) with manganate(VII) ions

Example 1:

Finding the mass of iron in an iron tablet: 


A typical multivitamin with iron tablet weighs 0.364g and is said to contain 14mg of iron. 

Is this the case?

To determine the iron content of this tablet first weigh two tablets and then grind them up and dissolve them in dilute (0.5M) sulphuric acid and make up the solution to 250ml in a volumetric flask. 

The role of the sulphuric acid is there to prevent oxidation of the iron (II) ions to iron(III) ions.  I’m here assuming that the iron in the tablet is in the form of iron(II)sulphate.   

Next titrate a 25ml aliquot of the solution from the volumetric flask with potassium manganate(VII) solution (0.0005M). 

Manganate(VII) ions oxidise the iron(II) ions according to the equation below.

5Fe²⁺      +      MnO₄^—    +        8H⁺               5Fe³⁺      +      Mn²⁺      +     4H₂O

The titration results show that it requires 20ml of the manganate(VII) solution to completely oxidise the iron(II) ions.

The end point of the titration is when the iron(II) solution turns a very pale purple because manganate(VII) is the purple and it is therefore self-indicating.

Calculation:

a)    Calculate moles of manganate(VII) ions from the titration. 

Use n=cV

Therefore n=  0.0005  ×   20/1000  =   0.00001 moles manganate(VII) ions.

b)    Determine the number of moles iron(II) ions in the aliquot from the equation since 5 moles iron(II) ions react with 1 moles manganate(VII) ions.

5Fe²⁺         +      MnO₄^—   

0.00005         0.00001

c)    if there are 0.00005 moles iron(II) ions in the aliquot then there are 10× this number of moles in the volumetric flask i.e. 0.0005 moles.

d)    Calculate the mass of iron that this number of moles represents.

Use m  =   n  ×  M  therefore mass Fe²⁺ ions =  0.0005  ×  55.9  =  0.02795g  or  27.95mg.

This mass of iron corresponds closely to the mass of iron reportedly present in 2 of these tablets 28mg. 

Example 2:

Finding the percentage of iron in a piece of steel weighing 0.034g.

First thing to do is dissolve the steel in the minimum volume of 2M sulphuric acid.  You may need to heat the flask in which you carry out this part of the experiment and do so in fume cupboard because of the hydrogen fumes that will be evolved. 

The resultant solution should look pale green in colour showing the presence of iron(II) ions. 

Next allow the solution to cool and then transfer it all to a 250ml volumetric flask and make up to the mark with distilled water. 

Next prepare a solution of potassium manganate(VII) 0.0005M and titrate a 25ml aliquot of the iron(II) ions solution with the purple manganate(VII) solution.

If the titration results show that 23.7ml of the manganate(VII) solution completely oxidise the iron(II) ions, calculate the % of iron in the sample of steel. 

a)    Calculate moles of manganate(VII) ions from the titration. 

Use n=cV

Therefore n=  0.0005  ×   23.7/1000  =   0.00001185 moles manganate(VII) ions.

b)    Determine the number of moles iron(II) ions in the aliquot from the equation since 5 moles iron(II) ions react with 1 moles manganate(VII) ions.

5Fe²⁺                +            MnO₄^—   

0.00005925              0.00001185

c)    if there are 0.00005925 moles iron(II) ions in the aliquot then there are 10× this number of moles in the volumetric flask i.e. 0.0005925 moles.

d)    Calculate the mass of iron that this number of moles represents.

Use m  =   n  ×  M  therefore mass Fe²⁺ ions =  0.0005925  ×  55.9  =  0.03312g.

e)    Therefore the percentage of iron = 0.03312/0.034  ×  100  =  97.4

Example 3:

Titration of iron(II) with ethandioate ions

Follow this link here to details about this titration.

9. Iron and Haemoglobin

Haem is an iron(II) complex with a multidentate ligand.

Oxygen forms a co-ordinate bond to Fe(II) in haemoglobin, enabling oxygen to be transported in the blood.

Carbon monoxide is toxic because it replaces oxygen co-ordinately bonded to Fe(II) in haemoglobin.

You can find out more about this area by following this link: Haemoglobin.

Transition metals: Some chemistry of iron(1)

Some chemistry of Iron(1):

Edexcel

24. be able to record observations and write suitable equations for the reactions of Cr³⁺(aq), Fe²⁺(aq), Fe³⁺(aq), Co²⁺(aq) and Cu²⁺(aq) with aqueous sodium hydroxide and aqueous ammonia, including in excess.

25. be able to write ionic equations to show the difference between ligand exchange and amphoteric behaviour for the reactions in (24) above.

34. understand the role of Fe²⁺ ions in catalysing the reaction between I− and S₂O₈^2— ions.

Manganate(VII) with iron (II) titration self-indicating

catalyst in Haber process

AQA

Exchange of the ligand H₂O by Cl– can involve a change of co-ordination number (e.g. Fe³⁺(aq), Co²⁺(aq) and Cu²⁺(aq).

Haem is an iron(II) complex with a multidentate ligand.

Oxygen forms a co-ordinate bond to Fe(II) in haemoglobin, enabling oxygen to be transported in the blood.

Carbon monoxide is toxic because it replaces oxygen co-ordinately bonded to Fe(II) in haemoglobin.

The redox titrations of Fe²⁺(aq) and C₂O₄^2– with MnO₄^—

Students should be able to perform calculations for these titrations and similar redox reactions.

Examples include, finding:

•   the mass of iron in an iron tablet 


•   the percentage of iron in steel 


•   the Mr of hydrated ammonium iron(II) sulfate 


•       Fe is used as a heterogeneous catalyst in the Haber process.

In aqueous solution, the following metal-aqua ions are formed:

[M(H₂O)₆]²⁺, limited to M = Fe and Cu

[M(H₂O)₆]³⁺, limited to M = Al and Fe

The acidity of [M(H₂O)₆]³⁺ is greater than that of [M(H₂O)₆]²⁺

Students should be able to:

•   explain, in terms of the charge/size ratio of the metal ion, why the acidity of [M(H₂O)₆]³⁺ is greater than that of [M(H₂O)₆]²⁺ 


•   describe and explain the simple test-tube reactions of M²⁺(aq) ions, limited to M = Fe and Cu, and of M³⁺(aq) ions, limited to M = Al and Fe, with the bases OH–, NH3 and CO₃^2— 


•       Students could carry out test-tube reactions of metal-aqua ions with NaOH, NH₃ and Na₂CO₃

Halogen carrier in arene organic chemistry

OCR

Redox reactions


(k) redox reactions and accompanying colour changes for:

(i) interconversions between Fe2+ and Fe3+

Fe2+ can be oxidised with H+/MnO4– and Fe3+ reduced with I–

Redox chemistry of iron and its compounds

1. Iron production in the Blast Furnace

Iron production is massive across the world.  It is still the most important piece of chemistry happening.  Nations across the world rely on iron and steel as construction materials for buildings and vehicles. 

And the material starts as iron(III) oxide Fe₂O₃

Fe₂O₃    +   3CO   ⟶     3CO₂   +   2Fe

Carbon monoxide generated in a blast furnace from the action of coke in oxygen–enriched air reduces the iron(III)oxide to iron which is then taken for processing into steel. 

. 

2. Relative reactivity of iron and copper

One of the most useful characteristics of iron is that it is not highly reactive and can be easily handled and stored as a solid unlike many more reactive metals like sodium which require protection on storage and handling. 

Furthermore, this low reactivity means that it does not corrode quickly and its corrosion can be easily prevented in several ways e.g. by painting or oiling or galvanising or using sacrificial protection.

Iron is easily plated if it is placed in a solution of copper ions. 

The redox chemistry is described in the equation below. 

CuSO₄   +    Fe     ⟶       FeSO₄     +   Cu

The reactivity series results from this difference in ability to react.

Highly reactive metals good at losing electrons from their outer shells we put top of the list and those metals unwilling to lose electrons are at the bottom. 

In the chart below, reactivity series is related to electrode potential.

3. Thermit reaction

This is the chemistry used to weld rail–track into continuous rail in situ.

2Al   +    Fe₂O₃      ⟶       2Fe     +     Al₂O₃

Aluminium powder reduces iron oxide to iron is the basic reaction but of course the metal mix has to match the steel qualities of the rails or else the weld would crack under use. 

5. Precipitation Reactions of iron(II) and iron(III).

With aqueous sodium hydroxide:  NaOH(aq)

Addition of sodium hydroxide solution to a solution of iron(I) or iron (III) ions produces distinctive precipitates both of which are insoluble in excess sodium hydroxide solution. 

Fe²⁺       +   2OH^—       ⟶         Fe(OH)₂(s)

                                                Pale green

The iron(II) hydroxide precipitate is a very pale green at first which turns darker green on standing due, actually, to oxidation if the iron(II) ions by dissolved oxygen in the water.  You can see in the photo below that the iron(II) hydroxide has darkened near the surface of the solution due to action of atmospheric oxygen.

Fe³⁺       +   3OH^—       ⟶         Fe(OH)₃(s)

                                                Dark brown

The iron(III) hydroxide precipitate is a dark rust brown in colour and it remains so on standing. 

These different colours serve to distinguish iron(II) from iron(III) ions. 

These precipitates are pictured below:

Similar reactions occur with alkaline aqueous ammonia (NH₃(aq))

6. Acidity of iron(III) ions:

The acidity of [Fe(H₂O)₆]³⁺ is greater than that of [Fe(H₂O)₆]²⁺

Iron(III) ions in aqueous solution are excellent proton donors with a significant pKa value. 

[Fe(H₂O)₆]³⁺       ⟶        [Fe(H₂O)₅ OH]²⁺      +      H⁺

pale lilac                               orange

Iron(II) ions by contrast are nowhere near as effective as proton donors and have a significantly different and larger pKa value. 

How can we explain the difference in acidity?

The difference in pKa values is down to the difference in charge/ size ratios or charge density values.

The iron(III) ion has a smaller radius and a greater charge giving it a greater charge density.

The greater the charge density (charge/size ratio) means that the iron(III) ion has greater polarising power acting on the water molecule ligands.

The effect is to distort the electron distribution in the water molecules and to therefore to weaken the OH bond in the water molecules and make the loss of a proton more likely. 

The diagram below shows how this works with large sodium and small iron(III) ions.

My next post will complete this brief survey of the chemistry of iron and some of its compounds.